Phosphorus Facts

Chemical & Physical Properties of Phosphorus

Pure phosphorus exists in several forms called allotropes.
Pure phosphorus exists in several forms called allotropes. This photo shows waxy white phosphorus (yellow cut), red phosphorus, violet phosphorus and black phosphorus. The allotropes of phosphorus have markedly different properties from each other. BXXXD, Tomihahndorf, Maksim, Materialscientist (Free Documentation License)

Phosphorus Basic Facts

Symbol: P

Atomic Weight: 30.973762

Discovery: Hennig Brand, 1669 (Germany)

Electron Configuration: [Ne] 3s2 3p3

Word Origin: Greek: phosphoros: light-bearing, also, the ancient name given the planet Venus before sunrise.

Properties: The melting point of phosphorus (white) is 44.1°C, boiling point (white) is 280°C, specific gravity (white) is 1.82, (red) 2.20, (black) 2.25-2.69, with a valence of 3 or 5. There are four allotropic forms of phosphorus: two forms of white (or yellow), red, and black (or violet). White phosphorus exhibits a and b modifications, with a transition temperature between the two forms at -3.8°C. Ordinary phosphorus is a waxy white solid. It is colorless and transparent in its pure form. Phosphorus is insoluble in water, but soluble in carbon disulfide. Phosphorus burns spontaneously in air to its pentoxide. It is highly poisonous, with a lethal dose of ~50 mg. White phosphorus should be stored under water and handled with forceps. It causes severe burns when in contact with skin. White phosphorus is converted to red phosphorus when exposed to sunlight or heated in its own vapor to 250°C. Unlike white phosphorus, red phosphorus does not phosphoresce in air, although it still requires careful handling.

Uses: Red phosphorus, which is relatively stable, is used to make safety matches, tracer bullets, incendiary devices, pesticides, pyrotechnic devices, and many other products. There is a high demand for phosphates for use as fertilizers. Phosphates are also used to make certain glasses (e.g., for sodium lamps). Trisodium phosphate is used as a cleaner, water softener, and scale/corrosion inhibitor. Bone ash (calcium phosphate) is used to make chinaware and to make monocalcium phosphate for baking powder. Phosphorus is used to make steels and phosphor bronze and is added to other alloys. There are many uses for organic phosphorus compounds. Phosphorus is an essential element in plant and animal cytoplasm. In humans, it is essential for proper skeletal and nervous system formation and function.

Element Classification: Non-Metal

Phosphorus Physical Data

Isotopes: Phosphorus has 22 known isotopes. P-31 is the only stable isotope.

Density (g/cc): 1.82 (white phosphorus)

Melting Point (K): 317.3

Boiling Point (K): 553

Appearance: white phosphorus is a waxy, phosphorescent solid

Atomic Radius (pm): 128

Atomic Volume (cc/mol): 17.0

Covalent Radius (pm): 106

Ionic Radius: 35 (+5e) 212 (-3e)

Specific Heat (@20°C J/g mol): 0.757

Fusion Heat (kJ/mol): 2.51

Evaporation Heat (kJ/mol): 49.8

Pauling Negativity Number: 2.19

First Ionizing Energy (kJ/mol): 1011.2

Oxidation States: 5, 3, -3

Lattice Structure: Cubic

Lattice Constant (Å): 7.170

CAS Registry Number: 7723-14-0

Phosphorus Trivia:

  • Hennig Brand isolated phosphorus from urine. He kept his process a secret, choosing instead to sell the process to other alchemists. His process became more widely known when it was sold to the French Academy of Sciences.
  • Brand's technique was replaced by Carl Wilhelm Scheele's method of extracting phosphorus from bones.
  • Phosphorus is the sixth most common element in the human body.
  • Phosphorus is the seventh most common element in the Earth's crust.
  • Phosphorus is the eighteenth most common element in seawater.
  • An early form of matches used white phosphorus in the match head. This practice gave rise to a painful and debilitating deformation of the jawbone known as 'phossy jaw' to workers when over-exposed to white phosphorus.

References: Los Alamos National Laboratory (2001), Crescent Chemical Company (2001), Lange's Handbook of Chemistry (1952), CRC Handbook of Chemistry & Physics (18th Ed.) International Atomic Energy Agency ENSDF database (Oct 2010)

Return to the Periodic Table